Redox Titrations with Permanganate and Dichromate

Introduction

Redox titrations use a reaction in which electrons move between analyte and titrant. Permanganate and dichromate are familiar oxidising titrants in acidic solution, but their reduction products and electron counts differ. A balanced half-reaction, not a colour or formula shortcut, determines the mole ratio. The acid conditions are part of the chemistry because changing pH can change the reduction product.

Core explanation

In sufficiently acidic solution, permanganate reduction is MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. Manganese moves from +7 to +2 and accepts five electrons per permanganate ion. Its intense purple colour can mark the endpoint: once reducing analyte is consumed, a slight excess of permanganate leaves a persistent pale pink tint. This is why permanganate is called self-indicating in many procedures. In neutral or alkaline media the product can differ, often including MnO₂, so the five-electron acidic equation must not be carried over blindly.

Dichromate reduction in acid is Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O. Two Cr(VI) centres each become Cr(III), accepting six electrons per dichromate. Dichromate's colour change is not typically used as a simple self-indicating endpoint in the same way as permanganate; a suitable redox indicator or instrumental detection may be used. Both titrants need appropriate standardisation and controlled acid conditions.

For Fe²⁺ analysis, oxidation is Fe²⁺ → Fe³⁺ + e⁻. One permanganate reacts with five Fe²⁺ in acidic conditions. One dichromate reacts with six Fe²⁺. For oxalate, C₂O₄²⁻ → 2CO₂ + 2e⁻, so two permanganate ions accept the ten electrons released by five oxalate ions. The overall reaction must also balance H⁺ and water. Some reactions, including permanganate with oxalate, can be kinetically slow at room temperature and require specified conditions; favourable electron bookkeeping alone does not guarantee a sharp practical endpoint.

The titration calculation follows n(titrant) = CV, then converts via the balanced electron ratio. If 1.00 mmol acidic permanganate is consumed entirely by Fe²⁺, it corresponds to 5.00 mmol Fe²⁺. Saying “one oxidant molecule reacts with one reductant molecule” without balancing electrons would yield a fivefold error.

Step-by-step reasoning

1. Determine oxidation-state changes for titrant and analyte under the specified pH conditions. 2. Balance oxidation and reduction half-reactions, including H⁺ and H₂O where appropriate. 3. Multiply half-reactions to cancel electrons and obtain the mole ratio. 4. Calculate titrant moles from standardised concentration and delivered volume. 5. Convert to analyte amount and check the endpoint signal and possible side reactions.

Visual explanation

Draw two electron ledgers. One MnO₄⁻ receives five electron tokens while five Fe²⁺ each supply one. One Cr₂O₇²⁻ receives six tokens while six Fe²⁺ supply one each. Beside the ledger draw a flask changing from colourless/pale reaction solution to a persistent faint permanganate tint after equivalence.

Real-world analogy

The titrant is an electron collector with a fixed number of empty slots. Acidic permanganate has five slots per ion and acidic dichromate six. The number of analyte particles needed depends on how many electrons each gives. Counting collector molecules without counting slots is like counting buses without knowing how many passengers each carries.

Real-world example

An analyst may titrate Fe²⁺ in an iron-containing solution using standardised permanganate under specified acidic conditions. Because atmospheric oxygen can oxidise Fe²⁺ during handling, sample preservation and rapid analysis matter. A correct five-to-one mole ratio cannot compensate for Fe²⁺ lost before titration.

Why?

Why is acid concentration specified in a permanganate method? Protons appear explicitly in the MnO₄⁻/Mn²⁺ half-reaction. If the medium is insufficiently acidic, a different manganese product may form and the electron ratio changes. The medium is therefore a reactant condition, not merely a background solvent choice.

Common misconception

“Permanganate always accepts five electrons” is only valid for reduction to Mn²⁺ under the stated acidic conditions. Another mistake is assuming dichromate is self-indicating just because it is coloured. Endpoint detection is a separate design choice from the oxidant's stoichiometry.

Worked example

An Fe²⁺ aliquot consumes 16.00 mL of 0.02000 mol L⁻¹ acidic KMnO₄. Permanganate moles = 0.01600 × 0.02000 = 0.0003200 mol. Each MnO₄⁻ accepts five electrons, and each Fe²⁺ supplies one, so Fe²⁺ moles = 5 × 0.0003200 = 0.001600 mol. If the aliquot was 25.00 mL, its Fe²⁺ concentration was 0.001600/0.02500 = 0.06400 mol L⁻¹, assuming no other reductant reacted.

Quick check

1. How many Fe²⁺ ions are oxidised per dichromate ion reduced to two Cr³⁺ ions in acid? Answer: Six Fe²⁺ ions, because dichromate accepts six electrons and each Fe²⁺ loses one when it becomes Fe³⁺.

Exam focus

Balance half-reactions under the actual pH conditions and explicitly count electrons. Distinguish equivalence from the observed endpoint. For permanganate, explain the persistent faint pink excess; for dichromate, state an appropriate separate endpoint method. Include any sample-preservation issue for oxidation-sensitive analytes.

Advanced insight

Redox potential changes with concentrations and pH according to the Nernst equation. A titration may be thermodynamically favourable yet show slow kinetics, interfering oxidation reactions or a gradual potential jump. The chosen acid and temperature are therefore parts of method validation rather than incidental details.

Summary

Acidic permanganate accepts five electrons per ion and can indicate its own slight excess by colour. Acidic dichromate accepts six electrons per ion and usually needs another endpoint signal. Balanced electron transfer fixes the analyte–titrant ratio; medium, kinetics and interferences determine practical reliability.

Practice questions

1. Write the acidic permanganate reduction half-reaction. Answer: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. It consumes five electrons per permanganate ion.

2. What moles of Fe²⁺ correspond to 0.00200 mol dichromate in acid? Answer: Six Fe²⁺ react per dichromate, so 0.0120 mol Fe²⁺ is oxidised to Fe³⁺.

3. Why might Fe²⁺ titration read low after a sample stands open to air? Answer: Dissolved oxygen may oxidise some Fe²⁺ to Fe³⁺ before the titration, leaving less Fe²⁺ to consume oxidising titrant.